Chapter 3

1. Chapter 3 Water and the Fitness of the Environment

2. Figure 3.1 Overview: The Molecule That Supports All of Life • Water is the most abundant biological medium here on Earth • All living organisms require water more than any other substance (50-95% of living organisms) • Three-quarters of the Earth’s surface is submerged in water • The abundance of water is the main reason the Earth is habitable

3. – Hydrogenbonds + H – + H + –  – + Figure 3.2 Concept 3.1: The polarity of water molecules results in hydrogen bonding • The water molecule is a polar molecule • Allows them to form hydrogen bonds with each other • hydrogen bonds are significantly weaker than ionic / covalent bonds (20 times weaker than covalent bonds ) • hydrogen bonds last only about 1/100,000,000,000th of a second, but hydrogen bonding between molecules is very important with organic compounds.

4. Hydrogen Bonds • Water molecule’s polarity is due to partial positive and partial negative ends. • +/- Polarity contributes to the various properties water exhibits • Each hydrogen of this water molecule can form hydrogen bonds with the oxygen atom of other water molecules = cohesion. • The hydrogen atoms of the water molecule can also form bonds with other slightly negative (polar) compounds = adhesion.

5. "Water is Wet" • Water adheres to a surface due to these two properties.

6. Concept 3.2: Four emergent properties of water contribute to Earth’s fitness for life Cohesion:Is the bonding of a high percentage of the molecules to neighboring water molecules Adhesion: The attraction between water and other substances. • These two properties allow for capillary action. • Water is attracted to the polar substances (adhesion) and climbs these substances, while pulling up the other water molecules due to cohesion. • The meniscus, in a column of water, is formed because gravity pulls down on the water molecules in the center while water molecules at the sides of the container "climb."

7. Water conducting cells 100 µm Figure 3.3 • Cohesion and adhesion help pull water up through the microscopic vessels (xylem) of plants “Imbibition” occurs when water moves into a substance (due to capillary action) and that substance swells.

8. Figure 3.4 Surface tension • Is a measure of how hard it is to break the surface of a liquid. • Water is attracted to itself, and this attraction, due to hydrogen bonds, is stronger than the attraction to the air above it. This causes molecules to hold together. • Water that is H-bonded together forms a “skin”.

9. Water’s High Specific Heat • The specific heat of a substance • Is the amount of heat that must be absorbed or lost for 1 gram of that substance to change its temperature by 1ºC • Water heats up as the hydrogen atoms vibrate (molecular kinetic energy = energy of molecular motion). • H2O has a high spec.heat because of many hydrogen bonds restricting movement of molecules. • trivia: H2O has 4x spec. heat of air (only ammonia is higher)

10. Moderation of Temperature • Water moderates air temperature, especially in coastal regions • Length of growing season • Types of species • Water’s high specific heat allows it to minimize temperature fluctuations to within limits that permit life. • In biological terms, the high spec. heat of H2O means that, for a given rate of heat input, the temperature of water will rise more slowly that the temperature of most other materials…Conversely, water will lose heat more slowly.

11. Evaporative Cooling • Evaporation, or “vaporization” • Is the transformation of a substance from a liquid to a gas • Heat of vaporization • Is the quantity of heat a liquid must absorb for 1 gram of it to be converted from a liquid to a gas • Water has a high heat of vaporization (over 500 cals per gram) because of the hydrogen bonds.

12. Evaporative cooling • Water is a good evaporative coolant, due to water’s high heat of vaporization • Allows water to cool a surface • Because it takes a lot of energy to change water from a liquid to a gas, when the vapor leaves it takes a lot of energy with it. • When humans sweat, water absorbs heat from the body. When the water turns into water vapor, it takes that energy (heat) with it.

13. Insulation of Bodies of Water by Floating Ice • Solid water, or ice • Is less dense than liquid water • Floats in liquid water • Water has a high freezing point and lower density as a solid than a liquid. Water's maximum density is 4ºC, while freezing solid is 0º C. • Water as a solid takes up more volume than liquid (ex; bursting water pipes) and is less dense (floats) than liquid (ex: frozen lakes). • This is why ice floats, this fact also allows for aeration of still ponds in spring and fall and the reason that ponds don't freeze from the bottom up.

14. Liquid water Hydrogen bonds constantly break and re-form Ice Hydrogen bonds are stable Hydrogen bond Figure 3.5 The hydrogen bonds in ice: • Are more “ordered” than in liquid water, making ice less dense • In most liquids, the density increases as temperature drops • In WATER, at just below 4 degrees Celsius, the individual molecules are so close that they form hydrogen bonds again which push them apart, creating a crystal lattice. Crystal lattice Since ice floats in water, life can exist under the frozen surfaces of lakes and polar seas

15. Heat of Fusion • "Heat of Fusion" 79.7 calories per gram to go from solid to liquid • H2O melting point 0º C / H2O freezing point 0º C • As ice melts, it takes in heat from the surroundings, cooling them down. • The presence of dissolved substances reduces the temperature at which water freezes or boils (ex: salt)

16. “The Solvent of Life” • Water is a versatile solvent due to its polarity • It can form aqueous solutions • hydrophobic (water-fearing) non polar substances tend not to dissolve in water. ex: lipids, petroleum, hydrocarbons • hydrophilic (water-loving) compounds; polar (charged), tend to dissolve in H2O • Water is able to dissolve anything polar due to its own polarity. • Water separates ionic substances +/- into their ions. • Many covalently bonded compounds may also have polar regions.

17. Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). – Na+ + + – + – – Positive hydrogen regions of water molecules cling to chloride anions (Cl–). Na+ – + + Cl – Cl– + – – + – + – – Figure 3.6 Water as a Solvent –summary: • Solution: uniform mixture of molecules of 2 or more substances • Solvent: (usually liquid) greatest in amount • Solute: (usually solid) lesser in amount The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them

18. Solute Concentration in Aqueous Solutions • Since most biochemical reactions occur in water, it is important to learn to calculate the [concentration] of solutes in an aqueous solution • A mole is the amount of an element equivalent to its atomic wt. in grams • A mole contains the same # of particles all the time: Avogadro’s #: 6.02 x 1023 • Molarity: is the number of moles of solute per liter of solution

19. + – H H H + H H H H H Hydroxide ion (OH–) Hydronium ion (H3O+) Figure on p. 53 of water dissociating Water Dissociates • …into H+ ions and hydroxide OH- ions • ideally, in pure water HOH H+ + OH- • In liquid H2O, the tendency exists for a hydrogen ion to "jump" onto other H2O molecules to which it is hydrogen bonded => H3O+ (called a HYDRONIUM ion)

20. Ionization of Water: Acids and Bases • Ideally, in pure water HOHH+ + OH- • When [H+] = [OH-], the pH is 7.0 (neutral) Acids: cause an increase in [H+], while lowering [OH-]Base: cause a decrease in [H+], while increasing [OH-] ex: HCl added, causes H+ > OH-ex: NaOH added, causes H+ < OH- The pH scale (0-14) measures degrees of acidity pH= negative log of the [conc.] of hydrogen ions in moles/liter          The lower the pH, the higher the [H+] (think inverse)

21. Concept 3.3: Dissociation of water molecules leads to acidic and basic conditions that affect living organisms

22. The Threat of Acid Precipitation • Acid Rain (precipitation) • Refers to rain, snow, or fog with a pH lower than pH 5.6 • Is caused primarily by the mixing of different pollutants with water in the air • Biggest culprit: coal-burning power plants • Sulfur dioxide, nitrous oxides

23. 0 1 Moreacidic 2 3 Acidrain 4 5 Normalrain 6 7 8 9 10 11 12 13 Morebasic 14 Figure 3.9 • Acid precipitation • Can damage life in Earth’s ecosystems

24. Strong and Weak Acids and Bases • Strong acids and bases ionize almost completely in water, causing an increase in H+ or OH- ions • Biologically Important Acids and Bases:carboxyl group: -COOH (acid)amino group: -NH2 (base)

25. pH Scale 0 1 Battery acid 2 Digestive (stomach) juice, lemon juice Vinegar, beer, wine, cola 3 Increasingly Acidic [H+] > [OH–] 4 Tomato juice 5 Black coffee Rainwater 6 Urine Neutral [H+] = [OH–] 7 Pure water Human blood 8 Seawater 9 10 Increasingly Basic [H+] < [OH–] Milk of magnesia 11 Household ammonia 12 Household bleach 13 Oven cleaner 14 Figure 3.8 • The pH scale and pH values of various aqueous solutions

26. Buffers • Are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution • Consist of an acid-base pair that reversibly combines with hydrogen ions • The internal pH of most living cells • Must remain close to pH 7

27. Buffers Help to create a constant pH by combining with H+ ions, removing or releasing them as pH falls (more H+) or rises (less H+), thus avoiding radical swings in pH Bicarbonate is the major buffer in human (blood) H2CO3 ------> HCO3- + H+ H2CO3 dissociates into H+ + HCO3- which can "soak up" excess acids or bases

28. Water on the web : http://wow.nrri.umn.edu/wow/index.html